The lithium market is divided into two distinct parts: lithium chemical and lithium mineral. Lithium produced as a chemical is used as a feedstock to produce other lithium compounds and metal. The vast range of final applications for lithium chemical include the manufacture of lithium batteries—primary and secondary—greases, glass and ceramics production, aluminum smelting, air quality control, catalysts, pharmaceuticals, polymers, cements, and alloys. As a mineral, lithium concentrates are used in the specialty glass and ceramics industry.
Recent estimates tentatively put the world consumption of lithium at close to 113,000 tonnes lithium carbonate equivalent (LCE) or 21,230 tonnes lithium metal in 2008. Details on total consumption and breakdown by end-use markets are seldom published because of the high degree of concentration within the lithium industry. The outlook for lithium consumption nonetheless appears optimistic with an overall growth rate of 5.8% predicted between 2008 and 2013. Demand for lithium carbonate, lithium hydroxide and lithium salts is projected to rise by 15% pa over the same period, from 3,940 t Li in 2008 to 7,720 t Li in 2013.
The mass production of plug-in hybrid and electric vehicles present the most significant upside potential for lithium demand in this end use. Estimates vary widely as to the market penetration these vehicles will achieve and as such lithium consumption could be significantly higher or lower than this 15% per year forecast. High fuel prices may force the industry toward formulating a “greener” environment, which in-turn could make lithium a strategic element causing an escalating growth in its demand.
SQM and Rockwood have both increased capacity at their operations in Chile and plan to increase further capacity in 2013. Capacity at the salt lakes in Tibet and China could also rise to 16,000 tonnes/year, although projects in both areas have had problems and capacity has thus far been underutilized and slow to ramp up. Several development projects are in various stages of planning worldwide, which may have an impact on the supply of lithium to the market.
The price for lithium chemicals and lithium minerals has consistently risen for the past several years. United States pricing for lithium carbonate ranges from $2.80-$3 per lb and $6,160-$6,600 per ton when sold as part of a large contract. There is no spot market for lithium and the price is negotiated on a contract basis.
Potassium chloride is the most popular potassium fertilizer, followed in a distant second place by potassium sulfate and then by potassium magnesium sulfate, potassium nitrate, potassium phosphate, and solutions of potassium thiosulfate and potassium polysulfide. The dominant use of potassium sulfate is as a source of the nutrients potassium and sulfur in high-value crops such as berries, citrus, spinach, lettuce, grapes, and tomatoes. It is also used to fertilize turf grass for golf courses and other landscaped and high-traffic grounds. potassium sulfate's industrial applications include the production of rubber, medicines, firebrick, and various construction materials.
More than 90% of the world's agricultural potassium requirements are supplied by potassium chloride, since it is a high-concentration fertilizer (60% K2O nutrient) that can be produced and supplied relatively cheaply per unit of K2O from a variety of sources. Global consumption of about 55 million tonnes of product has increased in response to growing populations and reduced arable land per capita requiring improved crop yield efficiencies. The remaining supply of nutrient potash is in the form of premium potash fertilizers, of which potassium sulfate is the predominate form, with relatively minor quantities of sulphate of potash magnesia and nitrate of potash being produced.
Statistics for potassium sulfate (containing approximately 50% K2O) consumption are uncertain since such a large percentage is accounted for by China, and even in the United States potassium sulfate may be folded in with statistics for potassium chloride. This market is estimated at 5 million tonnes worldwide. For the United States market, Great Salt Lake Minerals (“GLUM”), North America's leading producer of potassium sulfate from its solar evaporation facility at the Great Salt Lake in Utah, publishes details of its sales broken down into domestic and exports and reports United States Census Bureau numbers modified by the USGS. Based on these numbers (with the trade data presented in short tons), the annual apparent consumption peaked at more than 400,000 tons in 2005 and dropped to 329,000 tons in 2008 as higher prices reduced demand.
The United States regional breakdown of consumption is based on the size of the agricultural industry and the type of crops grown. In particular, California and Florida are important areas for cultivating fruit and vegetables and account for the bulk of potassium sulfate demand. Both states have a relatively high percentage of soil testing low to medium in potassium). More than 60% of United States fruit and nut production is in California, Oregon, and Washington, all state that use potassium sulfate. In addition, tobacco grown in the southeastern United States (North Carolina, Tennessee, Kentucky, and Virginia) is an important but declining consumer of potassium sulfate.
The world's potassium sulfate supply is derived from more than twenty producing companies, plus another 40 to 50 within China, each with an annual output ranging from less than 10,000 tonnes to more than 500,000 tonnes. Global capacity has grown to approximately 4.9 million tonnes/year, with the main producers being Asia (mainly China) (53%), Western Europe (29%), and North America (10%) followed by Latin America (4%, all Chile), the Former Soviet Union (3%), and the Middle East and Africa with less than 1% each. Based on the size of the market and the types of crops requiring potassium sulfate, the three most likely targets for potassium sulfate sales are: 1) California/Washington, 2) Florida and parts of the southeast coastline, and 3) the Kentucky/Tennessee region.
Statistics released by GSLM indicate that the average selling price of potassium sulfate FOB Ogden, Utah, for fiscal year (FY) 2002 and 2003 was about $210 per ton based on sales of about 250,000 tons. In FY 2004, prices increased to an average of $227 per ton as sales jumped more than 50% to 386,000 tons through GSLM's acquisition of IMC Global's customer base. GSLM purchased IMC's potassium sulfate capacity at Carlsbad, N. Mex., which it closed, but continued to service the customer base from its source in Utah. In 2005 and 2006, sales were steady and the average price increased to almost $260 and then to $292 per ton. In 2007, the volume of sales increased to 423,000 tons and the average price was almost $322 per ton. Although the volume sold fell to below 400,000 tons in 2008, the average selling price increased to almost $596 per ton. This increase in the selling price per ton is well illustrated in the quarterly results, which show an average price of $752 per ton for Q3 2008 and $975 per ton for Q4 2008. These price levels are supported by the average value of imports from Germany at $934 per ton in 2008. Other values based on significant quantities include Canada at $551 per ton.
A wide variety of processes have been developed to produce potassium sulfate and similar compounds. One example is the process practiced at the Arad facility in the Negev desert in Israel. Brine is taken from the Dead Sea and heated to high temperatures in a fluidized bed. The brine decomposes, releasing HCl among other compounds. The HCl is used to make hydrochloric acid. This in turn is reacted with mined phosphate rock to form potassium sulfate. See also U.S. Pat. No. 5,552,126 to Efriam et al. Potassium sulfate may be made directly from brines. One example is the process described in U.S. Pat. No. 3,977,835 to Chemtob et al. Readily processable salt groupings were selectively crystallized out of a complex salt brine from Sealres Lake containing potassium, sodium, chloride, sulfate, carbonate and borate ions by cooling the brine in at least one artificial cooling stage to a temperature sufficiently low to at least crystallize mirabilite, evaporating the brine in a first solar evaporator to crystallize out halite, or halite and burkeite, free of potassium salt values, and then further concentrating the brine in a second solar evaporator to obtain a grouping of salts rich in potassium values.
The brine contained potassium ion in an amount up to about 3% by weight, preferably from about 0.5% to about 2% by weight, carbonate ion in an amount of from about 2.5% to about 4.5% by weight, sulfate ion in an amount of from about 3.0% to about 6.0% by weight, and borate ion in an amount of from about 0.6% to about 1.2% by weight, all based on the total weight of the brine, with the balance of the ionic species present being sodium ion and chloride ion. This brine could be artificially cooled to temperatures as low as about −20° C. without crystallization of potassium salts.
After artificial cooling to at least 20° C. to crystallize at least mirabilite, the resultant brine was processed in a solar-evaporation stage to crystallize halite or a mixture of halite and burkeite, again, without crystallization of potassium salts. After solar evaporation to crystallize sodium salts, the brine was passed to another solar-evaporation stage where the potassium salts were deposited along with borax and sodium salts. Depending on the degree of artificial cooling, the relative amounts of glaserite and sylvite deposited varied. Sylvite was the most desired form and a high degree of artificial cooling was preferred. Selective salt group crystallization using cooling in combination with solar evaporation allowed a total harvesting of all salts contained in a complex brine.
A wide variety of processes have been developed to produce lithium carbonate and similar compositions. As with potassium sulfate, lithium carbonate has been recovered directly from brines. One such process is disclosed in U.S. Pat. No. 4,287,163 to Garrett et al., which involved use of soluble sulfate salts as salting-out agents to precipitate lithium sulfate monohydrate. Magnesium sulfate was a preferred salting-out agent. Other sulfate salts found useful as salting-out agents were sodium sulfate and sulfuric acid, including any of their hydrates (including magnesium hydrates) or partially dehydrated salts. Process solutions were concentrated in solar ponds.
Lithium carbonate and similar compositions also can be produced directly from ore that is rich in convertible lithium compounds. One example is the experimental process developed by the United States Bureau of Mines (“USBM”) in 1988. Lithium and Its Recovery from Low-grade Nevada Clays, USBM Bulletin 691 (1988). USBM worked with clays obtained from the McDermitt Caldera in Nevada, which contained lithium in the form of hectorite.
The experimental USBM process is depicted in FIG. 1. To convert the hectorite-bound lithium to lithium carbonate, the clay was mixed with limestone and gypsum and the mix was subjected to feed preparation 101 followed by roasting 102. The clay was soft and friable and required no heavy crushing. However, it was air-dried and passed through a jaw crusher to produce a minus 10-mesh material before blending. The limestone and gypsum were treated similarly. Further feed preparation 101 entailed grinding and mixing the ingredients for 1 hour in a ball mill. The resultant mixture (80% finer than 200 mesh) was pelletized with water to produce nominal 6.5 mm diameter pellets. These pellets contained up to 20% moisture and were dried at 700° C. before roasting step 102.
Objectives of roasting step 102 were to (1) generate calcined material for leaching 103, purification, and product recovery studies, (2) determine optimum roasting conditions, and (3) determine typical gas emissions. Batch tests were conducted in the roaster to determine optimum retention time and roast temperature. Small charges (500 g) of pelletized 5:3:3 mix were roasted. The test results showed a 2 hour retention time and 9000 C to be optimum. This retention time was used throughout the roaster studies; the temperature was varied in a few tests in which the effect of temperature on lithium extraction was investigated.
To generate calcine for use in product-recovery solution studies an equivalent 5:3:3 mixture of clay, limestone, and gypsum was used; batch testing had established the 5:3:3 mix as optimum. The pelletized feed was charged to the roaster 102 in 600-g increments every 5 minutes to simulate continuous operation. Generally, each test produced 80 lb. of calcine in 6.5 hour operating time.
The final phase of the roast studies involved investigating the effects of charge composition and roast temperature on lithium extraction. A test series was conducted in which various mixes were roasted. Lithium extraction was determined by water-leaching 103 composite samples of the calcines. Lithium extractions of at least 80% were attained with a wide range of clay-limestone-gypsum ratios. Also, good lithium extraction was achieved over a temperature range of 850° C. to 975° C. The 5:2:2 mix was chosen for cost evaluation because this mix provided high extraction with a relatively low reagent addition. Emissions of SO2 and fluorine were calculated from material balances. In a commercial operation, these off gases would require scrubbing before being vented to the atmosphere.
The objective of the leaching 103 tests was to determine the relationship between leach-system variables and optimal lithium extraction. The following variables were studied:
1. leach pulp percent solids,
2. wash water recycle,
3. calcine particle size, and
4. leach time.
The calcines leached in these leaching 103 tests were produced by roasting 5:3:3 mixtures of clay, limestone, and gypsum. Generally, 70 lb. of calcine was water-leached in each test. A slurry filter 104 recovered the leach solution. The filter cake from filter 104 was washed and discarded.
A series of 30-minute 103 leach tests was conducted at ambient temperature to study the effect of percent solids and wash water recycle on lithium extraction. The test results showed that the calcine was leached effectively at 40% solids with recycled wash water. At 50% solids, the lithium extraction decreased. Since the wash water was recycled to the leach step, the volume of wash water used was equal to the volume of makeup water required for the next leach.
The calcine pellets did not break apart during the leach step 103. USBM concluded that if a coarse particle could be leached effectively, grinding requirements would be minimized. A test series was conducted to study the effect of calcine particle size on lithium extraction. The calcine was leached for 30 minutes at 40% solids using recycled wash water. Test results showed that the 30-minute leach 103 extracted the lithium equally well from all particle sizes tested. To determine the effect of leach time on lithium extraction, a test series was conducted with coarse-crushed and whole pellets. The pellets were leached at 40% solids in recycled wash water.
Test results showed that lithium was extracted from coarse-crushed pellets with a 5-minute leach; whole pellets were not effectively leached in 5 minutes. Although the pellets did not break apart during the leach, prolonged agitation generated fines which affected filtration rates. Filtrate rates decreased with increased leach time and increased particle size. For 30-minute leaches, the whole pellet slurry filtered slowly because the filter cloth was blinded with fines. As the calcine particle size decreased, the fines tended to remain on top of the filter cake, allowing faster filtration. Overall, the test results indicate that coarse-grinding the calcine and 103 leaching it for 5 minutes at 40% solids provided good extraction and high filtration rates. Under these conditions, lithium extractions of 82 to 84% could be expected; the leach solution generally contained 2.5 to 3.0 g/l lithium.
USBM evaporator 105 was fed with leach solution and solution recycled from the previous test. The recycled solution—mother liquor plus product wash—accounted for about 20% of the total volume in evaporator 105. In addition to concentrating the solution, calcium as calcium carbonate was removed from the leach solution in this step of the process. The leach solution was saturated with calcium sulfate (about 0.6 g/l calcium ion). It was found that reducing the calcium ion concentration to about 0.015 g/l prevented calcium contamination of the product. The evaporation procedure involved the following steps:
1. The solution—leach plus recycle—was evaporated to about 50% of its original volume and then passed through filter 106 to remove calcium carbonate. Carbonate ion (approximately 15 g/l) present in the recycled solution precipitated over 99% of the calcium contained in the leach solution.
2. The filtrate was returned to evaporator 107. Evaporation continued until the solution was reduced to 20% of its original volume.
3. The hot concentrated solution, containing 12 to 13 g/l lithium, was transferred to the product precipitation step 108. Generally, this concentrated solution was cloudy because a small amount of lithium carbonate precipitated during evaporation.
Lithium recovery step 108 involved heating the concentrated solution to boiling and adding a stoichiometric amount of sodium carbonate to precipitate a lithium carbonate product. The objective of this step was to recover a product of at least 99% purity. Initially, the product was recovered from the hot solution by vacuum filtration and then dried. This procedure yielded a product of about 80% purity with the principal contaminants being sodium sulfate and potassium sulfate.
Numerous tests were conducted using leach solution to investigate product purification techniques. Test results were erratic because precise control of solution concentration was difficult and synthetic solutions were used to study operating variables. A series of laboratory tests was conducted using 1-liter batches of synthetic concentrated solution—made up with reagent chemicals—containing 97 g/l lithium sulfate, 158 g/l potassium sulfate, and 87 g/l sodium sulfate. Adding a stoichiometric amount of calcium carbonate to the hot solution precipitated lithium carbonate.
Product filtration and washing procedures were then studied. The following observations were made:
1. Pressure filtration 109 yielded a product of higher purity than vacuum filtration 109 by reducing the moisture content of the filter cake.
2. With pressure filtration 109, 4 to 6 liter of wash water per kilogram of dry product was required to produce a 99% pure product. A much higher volume of water was needed to produce a comparable product by vacuum filtration 109.
3. For pressure filtration 109, wash water volumes above 6 l/kg of dry product did not further improve product purity. Also, single-stage washing was as effective as either multistage washing or product reslurry.
4. Adding calcium carbonate as a saturated solution, rather than as a dry powder, had little effect on product purity. However, this procedure generated a coarse grainy product, in contrast to the fine powdery product obtained by adding dry calcium carbonate.
The wash water and product filtrate recovered in these tests contained 14 to 16 g/l of lithium carbonate. The wash was recycled to the evaporator. After a crystallization step, the mother liquor was also recycled.
In addition to residual lithium carbonate, the product filtrate from filter 109 contained high concentrations of potassium sulfate and sodium sulfate (over 150 g/l of each), preventing effective recycling. Tests showed that the most effective method for reducing the sulfate concentration involved crystallizing the salts in crystallizer 110 by chilling the product filtrate to between 0° and −4° C.; below −4° C. the filtrate froze. The mother liquor, which contained 70 g/l sodium sulfate and 100 g/l of potassium sulfate, was recovered by either vacuum or pressure filtration. Pressure filtration tended to reduce lithium loss by decreasing the amount of mother liquor present in the filter cake.
The filter cake was a mixture of glaserite and glauber salt Laboratory tests showed that glaserite and glauber salt could be recovered separately by a two-step crystallization procedure. At product filtrate temperatures down to about 17° C., the glaserite crystallized. The salt was recovered by vacuum filtration and analyzed as 33 wt % potassium, 8 wt % sodium, and <0.1 wt % lithium. Further cooling of the product filtrate (to as low as −4° C.) crystallized the glauber salt. The salt was recovered by pressure filtration and analyzed as 28 wt % sodium, 6 wt % potassium (a small amount of glaserite crystallized with the glauber salt), and 0.15 wt % lithium.
USBM roast-leach test results indicated 82% to 84% lithium extraction as optimum. Treating the leach solution by the methods specified resulted in 95% to 98% recovery of the contained lithium. Losses occurred in calcium carbonate filtration (0.5% loss) and in the crystallization step (2% to 5% loss, depending on the filtration method used to separate the mother liquor from the salts). Overall, 78 to 82% of the lithium contained in the clay was recovered as 99% pure lithium carbonate.
A material balance and a cost evaluation for a 5:3:3 ratio of clay-limestone-gypsum was prepared by USBM. The operating cost for this feed ratio was $2.12/1b lithium carbonate; this figure was revised to $2.27/lb. lithium carbonate as of May 1985. The cost evaluation showed raw materials (primarily limestone and gypsum) used in the process to be the most costly component. To lessen this expense, the evaluation recommended a reduction in the quantity of reagents used in the roast feed. The feed ratio could be reduced from 5:3:3 to 5:2:2 without affecting lithium extraction.
In May 1985, the USBM prepared a further cost evaluation. The evaluation estimated the operating cost of the process at $1.86/lb lithium carbonate produced. The evaluation was revised as of July 1987 and the updated operating cost was $2.02/lb lithium carbonate. The selling price of lithium carbonate as of July 1987 was $1.50/lb. As in the initial cost evaluation, a principal cost was the purchase of limestone and gypsum raw materials, which amounted to $0.39/lb. lithium carbonate. A high-cost section process section was evaporation, which, because of high fuel costs, added about $0.35/lb lithium carbonate produced. The capital cost was estimated to be about $105 million. No cost was allowed for land acquisition, mine development, and royalties on the ore. These costs would have to be considered before development of the resource could occur.
USBM concluded that for its process to be economical in the market of the time, the operating costs had to be reduced. The process unit operations that showed promise for cost reduction were identified as roasting 102, leaching 103, and evaporation 105/107. Limited rotary roaster testing was conducted to determine if the reagent requirement for the roast could be further reduced. The testing involved recycling lithium carbonate product filtrate and product wash water, without salt removal, to the roast step. Test results indicated that the salts in the recycled solution improved lithium recovery for the 5:1.5:1.5 feed ratio by about 1%; recovery for a 5:2:2 mix was improved by 3% to 5%.
In the USBM's proposed modification, the recycled solution would be used in pelletizing the roast feed; thus, the water requirements for feed preparation, as well as evaporation load, would be decreased. To reduce the capital costs associated with agitation leaching, percolation leaching was investigated. Tests were conducted using calcine from a 5:2:2 roast. The calcine was leached in a series of four acrylic columns measuring 4 inches in diameter by 4 feet high. Preliminary test results indicated that lithium extraction was comparable to that achieved using agitation leaching. The percolation leach solution contained about 7 g/l lithium; leach solution obtained by agitation leaching generally contained 2.5 g/l to 3 g/l lithium. This increase in solution loading could significantly reduce the evaporation load.
USBM believed that a cost-saving alternative for the evaporation step 105/107 would involve the use of solar evaporation. Although solar evaporation tests were not conducted, this modification was considered by USBM to be a viable alternative because of the hot, dry Nevada climate. The potential process modifications would all require expanded investigation to determine if they were viable alternatives to established procedures.